Chemistry Lab Notebook | Lab 4: Galvanic & Electrolytic Cells
Printed Name: Signature: Date:
_______________________ ___________________________ _____________________________
___
Name: Date: Experiment #: 4 Title: Galvanic &
________________ _________________ Electrolytic Cells
Purpose: To construct and measure four galvanic cells using different metal pairings, compare measured voltages against
theoretical E°cell values calculated from standard reduction potentials, and perform copper electroplating in an electrolytic
cell to compare theoretical mass gain (via Faraday's Law) against the experimentally measured mass gain.
Procedure Data / Results / Calculations
Part A: Galvanic Cells Standard Reduction Potentials Used:
1. Review the standard reduction potential table Cu²⁺/Cu = +0.34 V | H⁺/H₂ = 0.00 V | Sn²⁺/Sn = −0.1375
for Cu, H, Sn, Fe, and Zn. Determine which V
metal will act as cathode (higher E°, undergoes Fe²⁺/Fe = −0.447 V | Zn²⁺/Zn = −0.7618 V
reduction) and which as anode (lower E°,
undergoes oxidation) for each pairing.
Cell 1 — Copper / Zinc:
Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) E° = +0.34 V
2. Set up the galvanic cell using a porous cup
Anode: Zn(s) → Zn²⁺(aq) + 2e⁻ E° = −0.7618 V
nested inside a larger beaker to allow ion
migration between half-cells. Place each metal Overall: Cu²⁺(aq) + Zn(s) → Cu(s) + Zn²⁺(aq)
strip in its matching metal-ion solution. Connect Theoretical: 0.34 − (−0.7618) = 1.103 V | Measured:
the cathode wire (red) and anode wire (black) to 1.09 V
the multimeter.
Cell 2 — Copper / Iron:
3. For each of the four pairings, record the Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) E° = +0.34 V
measured voltage from the multimeter. Anode: Fe(s) → Fe²⁺(aq) + 2e⁻ E° = −0.447 V
Overall: Cu²⁺(aq) + Fe(s) → Cu(s) + Fe²⁺(aq)
4. Write out the two half-reactions (cathode as
Theoretical: 0.34 − (−0.447) = 0.787 V | Measured:
reduction, anode as oxidation). Cancel electrons
0.62 V
and combine into one balanced overall cell
equation. Calculate theoretical E°cell:
E°cell = E°cathode − E°anode Cell 3 — Copper / Tin:
Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) E° = +0.34 V
5. Compare theoretical E°cell to the measured Anode: Sn(s) → Sn²⁺(aq) + 2e⁻ E° = −0.1375 V
voltage for each trial. Note: a negative measured Overall: Cu²⁺(aq) + Sn(s) → Cu(s) + Sn²⁺(aq)
voltage indicates the electrodes are reversed Theoretical: 0.34 − (−0.1375) = 0.4775 V | Measured:
from the assumed configuration — swapping 0.50 V
them yields the correct positive voltage.
Cell 4 — Tin / Zinc:
Cells tested: Cathode: Sn²⁺(aq) + 2e⁻ → Sn(s) E° = −0.1375 V
• Copper – Zinc Anode: Zn(s) → Zn²⁺(aq) + 2e⁻ E° = −0.7618 V
• Copper – Iron Overall: Sn²⁺(aq) + Zn(s) → Sn(s) + Zn²⁺(aq)
• Copper – Tin Theoretical: −0.1375 − (−0.7618) = 0.6243 V |
• Tin – Zinc Measured: −0.58 V*
*Negative reading indicates electrodes were connected
in reverse. Swapping gives +0.58 V, consistent with the
theoretical 0.6243 V.
Summary:
Cell Cathode Anode Theo. E° Meas. E°
Cu–Zn Cu (+0.34) Zn 1.103 V 1.09 V
(−0.7618)
Cu–Fe Cu (+0.34) Fe 0.787 V 0.62 V
(−0.447)
Cu–Sn Cu (+0.34) Sn 0.4775 V 0.50 V
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Printed Name: Signature: Date:
_______________________ ___________________________ _____________________________
___
Name: Date: Experiment #: 4 Title: Galvanic &
________________ _________________ Electrolytic Cells
Purpose: To construct and measure four galvanic cells using different metal pairings, compare measured voltages against
theoretical E°cell values calculated from standard reduction potentials, and perform copper electroplating in an electrolytic
cell to compare theoretical mass gain (via Faraday's Law) against the experimentally measured mass gain.
Procedure Data / Results / Calculations
Part A: Galvanic Cells Standard Reduction Potentials Used:
1. Review the standard reduction potential table Cu²⁺/Cu = +0.34 V | H⁺/H₂ = 0.00 V | Sn²⁺/Sn = −0.1375
for Cu, H, Sn, Fe, and Zn. Determine which V
metal will act as cathode (higher E°, undergoes Fe²⁺/Fe = −0.447 V | Zn²⁺/Zn = −0.7618 V
reduction) and which as anode (lower E°,
undergoes oxidation) for each pairing.
Cell 1 — Copper / Zinc:
Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) E° = +0.34 V
2. Set up the galvanic cell using a porous cup
Anode: Zn(s) → Zn²⁺(aq) + 2e⁻ E° = −0.7618 V
nested inside a larger beaker to allow ion
migration between half-cells. Place each metal Overall: Cu²⁺(aq) + Zn(s) → Cu(s) + Zn²⁺(aq)
strip in its matching metal-ion solution. Connect Theoretical: 0.34 − (−0.7618) = 1.103 V | Measured:
the cathode wire (red) and anode wire (black) to 1.09 V
the multimeter.
Cell 2 — Copper / Iron:
3. For each of the four pairings, record the Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) E° = +0.34 V
measured voltage from the multimeter. Anode: Fe(s) → Fe²⁺(aq) + 2e⁻ E° = −0.447 V
Overall: Cu²⁺(aq) + Fe(s) → Cu(s) + Fe²⁺(aq)
4. Write out the two half-reactions (cathode as
Theoretical: 0.34 − (−0.447) = 0.787 V | Measured:
reduction, anode as oxidation). Cancel electrons
0.62 V
and combine into one balanced overall cell
equation. Calculate theoretical E°cell:
E°cell = E°cathode − E°anode Cell 3 — Copper / Tin:
Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) E° = +0.34 V
5. Compare theoretical E°cell to the measured Anode: Sn(s) → Sn²⁺(aq) + 2e⁻ E° = −0.1375 V
voltage for each trial. Note: a negative measured Overall: Cu²⁺(aq) + Sn(s) → Cu(s) + Sn²⁺(aq)
voltage indicates the electrodes are reversed Theoretical: 0.34 − (−0.1375) = 0.4775 V | Measured:
from the assumed configuration — swapping 0.50 V
them yields the correct positive voltage.
Cell 4 — Tin / Zinc:
Cells tested: Cathode: Sn²⁺(aq) + 2e⁻ → Sn(s) E° = −0.1375 V
• Copper – Zinc Anode: Zn(s) → Zn²⁺(aq) + 2e⁻ E° = −0.7618 V
• Copper – Iron Overall: Sn²⁺(aq) + Zn(s) → Sn(s) + Zn²⁺(aq)
• Copper – Tin Theoretical: −0.1375 − (−0.7618) = 0.6243 V |
• Tin – Zinc Measured: −0.58 V*
*Negative reading indicates electrodes were connected
in reverse. Swapping gives +0.58 V, consistent with the
theoretical 0.6243 V.
Summary:
Cell Cathode Anode Theo. E° Meas. E°
Cu–Zn Cu (+0.34) Zn 1.103 V 1.09 V
(−0.7618)
Cu–Fe Cu (+0.34) Fe 0.787 V 0.62 V
(−0.447)
Cu–Sn Cu (+0.34) Sn 0.4775 V 0.50 V
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